Acid–Base Theories, Buffers, Indicators, Titrations & pKa List

Acid–Base Theories

Arrhenius (1883)

Acid: A substance that yields H⁺ in aqueous solution. Base: A substance that in aqueous solution gives OH^-.

Limitations: substances with basic properties that do not contain hydroxyl ions (e.g., NH₃); limited to aqueous solutions. Free, isolated protons do not exist in water.

Brønsted–Lowry (1923)

Acid: A species that tends to donate an H⁺. Base: A species that tends to accept an H⁺.

Lewis (1923)

Acid: A species that can accept an electron pair. Base: A species that can donate an electron pair.

Buffer Solutions

1. Buffer (dissolution buffer)

Definition: Buffer solutions are solutions that maintain a nearly constant pH when small amounts of acid or base are added, or when diluted.

Composition: Significant amounts of a weak acid and its conjugate base (or a weak base and its conjugate acid).

Buffer capacity: The number of moles of acid or base that can be added to a buffer before the pH begins to change appreciably.

Indicators

2. Indicators

Indicators are weak acids or bases whose acid and conjugate base forms have different colors. Using the indicator ratio [HIn]/[In^-]:

• If [HIn]/[In^-] >= 10 → Color A (acid form predominates).
• If [HIn]/[In^-] <= 0.1 → Color B (base form predominates).
• If the ratio is between 0.1 and 10 → a mixture of colors (transition range).

Titrations

3. Titrations

A titration is a technique in which a solution of known concentration (the titrant) reacts with a solution of unknown concentration. We measure the volume of titrant required to completely neutralize the analyte; the point at which stoichiometric equivalence is reached is the equivalence point.

pH at the equivalence point depends on the strengths of the acid and base involved:

• Strong acid + strong base → pH = 7 at equivalence.
• Weak acid + strong base → pH > 7 at equivalence.
• Weak base + strong acid → pH < 7 at equivalence.

Common Acids and Ka Values

List of acids with approximate acid dissociation constants (Ka):

• Acetic — HC₂H₃O₂ — Ka = 1.8 × 10^-5
• Arsenic (arsenic acid) — H₃AsO₄ — Ka ≈ 5.6 × 10^-3
• Benzoic — HC₇H₅O₂ — Ka = 6.5 × 10^-5
• Butanoic (butyric) — HC₄H₇O₂ — Ka = 1.5 × 10^-5
• Carbonic (H₂CO₃) — Ka ≈ 4.3 × 10^-7
• Hydrocyanic (hydrogen cyanide) — HCN — Ka = 4.9 × 10^-10
• Citric — H₃C₆H₅O₇ — Ka = 7.4 × 10^-4
• Chloroacetic — HC₂H₂O₂Cl — Ka = 1.4 × 10^-3
• Phenol — HC₆H₅O — Ka = 1.3 × 10^-10
• Hydrofluoric — HF — Ka = 6.8 × 10^-4
• Formic — HCHO₂ — Ka = 1.8 × 10^-4
• Phosphoric — H₃PO₄ — Ka ≈ 7.5 × 10^-3
• Hypobromous — HBrO — Ka = 2.5 × 10^-9
• Hypochlorous — HClO — Ka = 3.0 × 10^-8
• Lactic — HC₃H₅O₃ — Ka = 1.4 × 10^-4
• Nitrous — HNO₂ — Ka = 4.5 × 10^-4
• Oxalic — H₂C₂O₄ — Ka = 5.9 × 10^-2
• Sulfide (hydrogen sulfide) — H₂S — Ka = 9.5 × 10^-8
• Sulfuric — H₂SO₄ — strong acid!
• Bisulfite (HSO₃^-) — Ka ≈ 1.7 × 10^-2
• Iodic — HIO₃ — Ka = 1.7 × 10^-1

Common Bases and Kb Values

List of bases with approximate base dissociation constants (Kb):

• Ammonia — NH₃ — Kb = 1.8 × 10^-5
• Aniline — C₆H₅NH₂ — Kb = 4.3 × 10^-10
• Ethylamine — C₂H₅NH₂ — Kb = 6.4 × 10^-4
• Methylamine — CH₃NH₂ — Kb = 4.4 × 10^-4

Note: All numerical values are approximate Ka or Kb values given for reference. Chemical names and formulas have been corrected for spelling and formatting while preserving the original content.